UNIT 5: *STUDENT*
MOLES &
STOICHIOMETRY
*STUDENT*
VOCABULARY:
1.
2.
3.
4.
5.
6.
7.
8.
Mole
Formula mass (FM)
Gram formula mass (GFM)
Coefficient
Subscript
Species
Law of conservation of mass
Law of conservation of energy
9. Balanced equation
10. Synthesis reaction
11. Decomposition reaction
12. Single-replacement reaction
13. Double-replacement reaction
14. Molecular formula
15. Empirical formula
16. Percent mass
1
INTRODUCTION:
Before we can even begin to understand what this unit is about, we need to be
able to find the mass of different compounds. Open your Periodic Table and we’ll
get started…
 First, what are the units we use for the mass atoms?
___________________________ (_____)
 What is the mass of one atom of oxygen?
_________________
 Why don’t we use grams as the units for massing atoms? Atoms
are too SMALL—the number would be VERY BULKY
Ex: If we used grams to mass atoms, the mass of oxygen would be
0.00000000000000000000027 g or 2.7 x 10-23g
Find the mass of the following atoms:
1) Mg = __________
3) Cl = __________
5) Ca = __________
2) Li = __________
4) Al = __________
6) H = __________
 ____________ ELEMENTS = one atom of an element that’s stable
enough to stand on its own (VERY RARE)—not bonded to anything
 ____________ ELEMENTS or DIATOMS = elements whose atoms
always travel in pairs (N2, O2, F2, Cl2, Br2, I2, At2, H2)—bonded to
another atom of the same element
So, what would the mass be of one molecule of oxygen (O2)?
O2
subscript = tells you the
total number of atoms in
the compound/molecule
This means that the mass of O2 = 2 x ______ amu = ________ amu
2
Calculating the Formula Mass & Gram Formula Mass of
Compounds:
 __________________: the mass of an atom, molecule or compound in
ATOMIC MASS UNITS (amu)
Ex: formula mass of a hydrogen atom is 1.00794 amu
 ___________________: the mass of one _______ of an atom,
molecule or compound in GRAMS (g)
Ex: GFM of hydrogen is 1.00794 g (this is the mass of 1 mol H)
 __________: 6.02 x 1023 units of a substance (like a really big dozen)
Ex: 1 mol of C = 6.02 x 1023 atoms of C = 12.0111 g of C
Practice – SHOW ALL WORK!
1) What is the formula mass of K2CO3?
2) What is the gram formula mass of CuSO45H2O?
3
Calculating Percent Composition
Step 1: Calculate the GFM for the compound (or the FM).
Ex: CaCl2
Ca = 1 x 40.08 =
Cl = 2 x 35.453 = ___________
(this is the “part” Ca)
(this is the “part” Cl)
Step 2: Check the last page of your periodic table for the formula for
percent composition. Write the formula below:
% composition by mass = mass of part X 100
mass of whole
Now, use the formula to find the percent composition of each element or
“part” in our compound (to the nearest tenth of a %).
4
Practice:
1) What is the percentage by mass of carbon in CO2?
2) What is the percent by mass of nitrogen in NH4NO3?
3) What is the percent by mass of oxygen in magnesium oxide?
5
A Special Type of Percentage Composition—HYDRATES:
A _____________ is a _______________ compound in which ions are
attached to one or more __________ molecules
Example: Na2CO310H2O
 Notice how __________ molecules are BUILT INTO the
chemical formula
 Substances without water built into the formula are called
___________________
Problem: What is the percentage by mass of water in sodium carbonate
crystals (Na2CO310H2O)?
Step 1: Calculate the formula mass for the hydrate.
Na =
2 x
=
C=
1 x
=
O=
3 x
=
H2 O =
10 x
=
Formula mass of hydrate
=
Step 2: Check the last page of your periodic table for the formula for
percent composition. Write the formula below:
% composition by mass = mass of part X 100
mass of whole
% H2O by mass =
6
Practice:
1) What is the percent by mass of water in BaCl22H2O?
2) Which species contains the greatest percent by mass of oxygen?
b) H2O
c) NO2
d) MgO
a) CO2
CO2
H2 O
NO2
MgO
7
Application of the Mole: The Math of Chemistry
We will need to convert from grams to moles and vice versa for this class.
The diagram below summarizes these processes:
Mass in
grams (g)
Divide by GFM
Multiply by GFM
Number of
moles
(mol)
Converting Grams to Moles:
From Table T, you would use the Mole Calculations Formula:
# of moles = given mass (g)
GFM (g/mol)
8
Problem: How many moles are in 4.75 g of sodium hydroxide? (NaOH)
Step 1: Calculate the GFM for the compound.
Na = 1 x
O =1x
H =1x
=
=
=
_
Step 2: Plug the given value and the GFM into the “mole calculations”
formula and solve for the number of moles.
# of moles = given mass (g) = ____________ =
GFM (g/mol)
Practice:
1) How many moles are in 39.0 grams of LiF?
2) What is the number of moles of potassium chloride present in 148 g?
3) How many moles are in 168 g of KOH?
9
Converting from Moles to Grams:
From Table T, you would still use the Mole Calculations Formula, but
you must rearrange it since you are solving for _________ now:
mass of sample (g) = # of moles (mol) x GFM (g/mol)
Problem: You have a 2.50 mole sample of sulfuric acid. What is the
mass of your sample in grams? (H2SO4)
Step 1: Calculate the GFM for the compound.
Step 2: Plug the given value and the GFM into the “mole calculations”
formula and solve for the mass of the sample.
mass of sample (g) = # of moles (mol) x GFM (g/mol)
Practice:
1) What is the mass of 4.5 moles of KOH?
10
2) What is the mass of 0.50 mol of CuSO4?
3) What is the mass of 1.50 mole of nitrogen gas?
4) What is the total mass of 0.75 moles of SO2?
11
Chemical Equations:
 A CHEMICAL EQUATION is a set symbols that state the _________
and ____________ in a chemical reaction.
_____________ = the starting substances in a chemical reaction
(found to the _______ of the arrow)
_____________ = a substance produced by a chemical reaction
(found to the _______ of the arrow)
Example:
2Na + 2H2O  2NaOH + H2
 Chemical equations must be __________. Think of the arrow ()
as an equal sign.
 LAW of CONSERVATION of MASS: mass can neither be __________
nor _______________ in a chemical reaction
Balancing Equations:
The number of _______ of each __________ on the __________ (left)
side of the equation must be the same as the number of _______ of each
___________ on the ______________ (right) side of the equation.
*_______________ and ______________ tell us how many
moles we have for each element
Let’s look at the BALANCED equation below:
C
+
O2

CO2
*Note that there is 1 mol of carbon and 2 mol of oxygen on each
side of the arrow. That’s what it means to be BALANCED.
12
Now, let’s examine the following UNBALANCED equation:
H2
+
O2
H2O

Q: How does this unbalanced equation violate the Law of Conservation
of Mass?
A: In this equation, oxygen would have to be ___________ (there’s one
less on the products side)
 ______________ = the integer in front of an element or compound
which indicates the number of moles present
 ____________ = the integer to the lower right of an element which
indicates the number of atoms present
 ___________ = the individual reactants and products in a chemical
reaction.
Q: What do we use to balance equations?
A: __________________
NOTE: WE NEVER CHANGE THE SUBSCRIPTS IN A
FORMULA!
Example:
2Ag + S

Ag2S
SUBSCRIPTS:
Ag = ___
S = ___
Ag2S:
Ag = ___
S = ___
COEFFICIENTS:
Ag = ___
S = ___
Ag2S = ___
13
Method for Balancing Equations:
Step 1: Draw a line to separate products from reactants
Step 2: List each of the different elements on each side of the line
Step 3: Count up the number of atoms on each side & record next to
the element symbol
Step 4: Find the most complex compound in the equation. Balance the
elements found in that compound on the opposite side of the
arrow by changing the coefficients for the different species.
Every time you change a coefficient, you must update the
number of each element.
Step 5: Now, continue balancing the elements by changing coefficients
until you have the same number of each element on both sides
of the equation.
Example:
____ H2
+
____ O2

____ H2O
Example:
____ Na + ____ H2O  ____ NaOH + ____ H2
Example:
____ CO2 + ____ H2O  ____ C6H12O6 + ____ O2
14
ONE LAST NOTE:
When balancing chemical equations, ______________________ may be
balanced as a _________________ rather than as separate elements as
long as they stay intact during the reaction.
Example:
Al2(SO4)3
+
Ca(OH)2

Al(OH)3
+
CaSO4
In this equation, we have the polyatomic ions SULFATE & HYDROXIDE,
and both remain intact during the reaction. Since SO4 has the subscript
of 3, we could think of it as 3 x 1 = 3 sulfur atoms and 3 x 4 = 12 oxygen
atoms. OR, we can just look at the UNIT and say there are 3 (SO4)’s on
the reactant side and 1 (SO4) on the product side.
Now let’s balance the equation:
___ Al2(SO4)3 + ___ Ca(OH)2
15

___ Al(OH)3 + ___ CaSO4
TYPES OF CHEMICAL REACTIONS:
Type 1: SINGLE REPLACEMENT
Type 2: DOUBLE REPLACEMENT
Definition: Reaction where one
species replaces another (one
species alone on one side and
combined on the other).
Definition: Reaction where
compounds react, switch partners
and produce 2 new compounds.
Ex:
3Ag + AuCl3 → 3AgCl + Au
2Cr + 3H2SO4 → Cr2(SO4)3 + 3H2
2Cr + 3FeCO3 → Cr2(CO3)3 + 3Fe
Ex:
Pb(NO3 )2 + 2NaCl → PbCl2 + 2NaNO3
Na3PO4 + 3AgNO3 → Ag3PO4 + 3NaNO3
K2CO3 + 2AgNO3 → Ag2CO3 + 2KNO3
Will look like:
Will look like:
____________________
____________________
Type 3: SYNTHESIS
Type 4: DECOMPOSITION
Definition: Reaction where we take
more than one reactant and create
one product
Definition: Reaction where we take
one reactant and create 2 products.
Ex:
4Al + 302 → 2Al2O3
2H2 + O2 → 2H2O
Ex:
BaCO3 → BaO + CO2
2H2O2 → 2H2O + O2
2Bi(OH)3 → Bi2O3 + 3H2O
Will look like:
Will look like:
____________________
____________________
16
Mole-Mole Problems: An Introduction
A chemical equation is basically the “recipe” for a reaction. The
______________ in an equation tell us the amounts of ____________ and
_____________ we need to make the recipe work. Reactants in an equation
react in specific _________ to produce a specific amount of products.
Below is a recipe for sugar cookies:
3 eggs + 1 cup of flour + 2 cups sugar → 24 cookies
Let’s simplify this to: 3E + 1F + 2S → 24C
If we massed the eggs, flour and sugar, they should (in a perfect world)
equal the mass of the cookies. (This illustrates the LAW OF
CONSERVATION OF MASS)
3E
(200 g
+ 1F + 2S = 24C
+ 160g + 240g = 600g)
Q: If you had to bake 48 cookies, how many eggs would you need?
A: ___ (double it)
 Method for solving mole-mole problems: set up a proportion using
your known and unknown values, then cross-multiply and solve for your
unknown.
Ex 1: Set up the proportion from the Q & A above and solve.
3E + 1F + 2S → 24C
48
=
24
3
____x = _____
x = __________
17
Ex 2: If you have 10 eggs and an infinite amount of sugar and
flour, what is the greatest number of cookies you can make?
10
=
24
3
___x = _____
x = _____________
*We can use the process we used with the cookie recipe and apply it to
chemical equations. The only difference is we ALWAYS check to make
sure we are starting with a BALANCED CHEMICAL EQUATION
Ex 3: Consider the following formula:
N2 + 3H2 → 2NH3
How many moles of nitrogen gas (N2) would be needed to produce
10 moles of ammonia (NH3)?
18
Mole-Mole Practice
Use the following equation to answer questions 1-3:
C3H8 + 5O2 → 3CO2 + 4H2O
1) If 12 moles of C3H8 react completely, how many moles of H2O are
formed?
2) If 20 moles of CO2 are formed, how many moles of O2 reacted?
3) If 8 moles of O2 react completely, how many moles of H2O are
formed?
Use the following equation to answer questions 4-7:
N2 + 3H2 → 2NH3
4) If 2.5 moles of N2 react completely, how many moles of NH3 are
formed?
5) If 9 moles of NH3 are formed, how many moles of H2 reacted?
6) If 3.5 moles of NH3 are formed, how many moles of N2 reacted?
7) How many grams of N2 are reacted when 3.5 moles of NH3 are
formed?
19
Determining EMPIRICAL Formulas:
Empirical Formula = the reduced formula; a formula whose subscripts cannot
be reduced any further
Molecular Formula = the actual formula for a compound; subscripts represent
actual quantity of atoms present
Molecular Formula
N2O4
C3H9
C6H12O6
B4H10
C5H12
Empirical Formula
Practice: Determining empirical formula from molecular formula.
a) H2O
h) C2H6
b) N2O2
i) Na2SO4
c) C3H8
j) C6H5N
d) Fe(CO)3
k) P2O5
e) C5H10
l) H2O2
f) NH3
m) SeO3
g) CaBr2
n) LiCl
20
Calculating Empirical Formula from % Mass
Step 1: Always assume you have a 100 g sample (The total % for the
compound must = 100, so we can just change the units from % to g)
Step 2: Convert grams to moles.
Step 3: Divide all mole numbers by the smallest mole number.
Ex. A compound is 46.2 % mass carbon and 53.8 % mass nitrogen.
What is its empirical formula?
Step 1: Assume a 100 g sample.
46.2 % C = 46.2 g C
53.8 % N = 53.8 g N
Step 2: Convert grams to moles (have grams, need moles)
But we must have WHOLE NUMBERS for SUBSCRIPTS.
Step 3: Divide each mole number by the smallest mole number (We
will round in this step to the nearest integer if it’s super close).
For C: =
For N: =
So, the empirical formula for our compound is _______
21
Practice: Determine the empirical formula from the percent
composition for each of the following.
1) A compound contains 24.0 g C and 32.0 g O. Calculate its empirical
formula. (Hint: start with step 2)
2) A compound contains 0.50 moles of carbon for each 1.0 mole of
hydrogen. Calculate the empirical formula of this compound.
(Hint: start with step 3)
3) A compound contains 14.6% C and 85.4% Cl by mass. Calculate the
empirical formula of this compound.
4) 32.8% chromium and 67.2% chlorine.
5) 67.1% zinc and the rest is oxygen.
22
Determining MOLECULAR Formulas:
So far, we know how to:
1. Find an empirical formula from percent mass
2. Find an empirical formula from a molecular formula
But how do we find out the molecular formula from an empirical
formula?
Ex: A compound is 80.0 % C and 20.0 % H by mass. If its molecular
mass is 75.0 g, what is its empirical formula? What is its molecular
formula?
First, we must determine the empirical formula using the 3-step
process.
Step 1: Assume a 100 g sample.
80.0 % C =
20.0 % H =
Step 2: Convert grams to moles (have grams, need moles)
Step 3: Divide each mole number by the smallest mole number and
round to the nearest integer
For C: =
For H: =
So, the empirical formula for our compound is ________.
23
Now we




can determine the molecular formula:
Empirical mass (the mass of 1 mole of CH3) = ______ g
Molecular mass = ______ g
Molecular mass is ___ times larger than empirical mass
Molecular formula must be ___ times larger than empirical
formula
 Multiply ALL the subscripts in our empirical formula by ___
Thus, our molecular formula is _________
Practice: Answer the questions below in the space provided. SHOW
ALL WORK.
1) What is the molecular formula of a compound that has an
empirical formula of NO2 and molecular mass of 92.0 g?
2) A compound is 50% sulfur and 50% oxygen by mass. Calculate the
empirical formula. If its molecular mass is 128 g, determine its
molecular formula.
3) A compound is 63.6% N and 36.4% O by mass. Calculate its
empirical formula. List three possible molecular formulas for this
compound.
24
4) A compound is 92.3% carbon and 7.7% hydrogen by mass.
Calculate its empirical formula. If the molecular mass is 78.0 g,
determine its molecular formula.
5) A compound is 74.0% C, 8.7% H, and 17.3% N. Calculate its
empirical formula. Its molecular mass is 162 g. Determine its
molecular formula.
25